I read in a book that bond dissociation energy of carbon monoxide $(\ce{CO})$ is $\pu{749 kJ/mol}$ and carbon dioxide $(\ce{CO2})$ is $\pu{532 kJ/mol}.$ Shouldn't it take twice as much energy to break $\ce{CO2}?$


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    $\begingroup$ 532 kJ/mol times two... $\endgroup$ – Karsten Theis Jun 12 at 14:47
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    $\begingroup$ I'm voting to close this question as off-topic because your question lacks context. This question needs revision before it is ready for a great answer. Please edit to include how this question came up and how you tried to answer it. This will help writing an answer that is useful for you and for others. $\endgroup$ – Karsten Theis Jun 12 at 18:20

The bonds are not equivalent, therefore they have different bond energy.

Carbon monoxide consists of one carbon atom and one oxygen atom, connected by a triple bond that consists of two covalent bonds as well as one dative covalent bond.It is the simplest oxocarbon and is isoelectronic with other triply-bonded diatomic molecules having ten valence electrons, including the cyanide anion, the nitrosonium cation and molecular nitrogen. In coordination complexes the carbon monoxide ligand is called carbonyl.

Source: Wikipedia page on Carbon Monoxide


The bond dissociation refers to each bond. Therefore, the number you cited for $\ce{CO2}$ is for only one of the C=O bonds.


Look at it simply. Bond strength is proportional to the number of bonds present. CO has a triple bond which is consisted by 1 sigma and 2 pi bonds. This kind of covalent bond is relatively hard to break compared to a single and double bond.

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